Final answer:
Zero-order reactions have a constant rate regardless of reactant concentration, while the rate of first-order reactions is directly proportional to one reactant's concentration. The half-life of zero-order reactions depends on initial concentration, unlike first-order reactions. A change in reaction order might reflect a reaction mechanism's complexity or external factors affecting reactant collisions.
Step-by-step explanation:
Understanding Zero-Order and First-Order Reactions
A zero-order reaction is characterized by a rate that is independent of the concentration of the reactants. This means that no matter the amount of reactant present, the rate at which the reaction occurs remains constant, as the rate law can be written with the exponent of the reactant being 0. The half-life of a zero-order reaction varies depending on the initial concentration, which contrasts with the half-life of a first-order reaction, that remains constant regardless of concentration. In the event where a reaction changes from zero-order to first-order when the concentration is halved, it suggests a complexity in the reaction mechanism that cannot be explained using simple kinetics. The integrated rate law for a zero-order reaction is given by [A] = -kt + [A]o, where [A] is the concentration of the reactant at time t, k is the rate constant, and [A]o is the initial concentration.
The Arrhenius equation and collision theory contribute to our understanding of these reactions by describing how temperature, activation energy, and the orientation of collisions impact reaction rates. The change in reaction order could be due to a shift in the rate-determining step or other factors that influence the collision frequency and energy of reactants.